AQA GCSE Chemistry 8462: The rate and extent of chemical change – Study Guide

This guide teaches section 4.6 The rate and extent of chemical change of the AQA GCSE Chemistry (8462) specification, for teaching from September 2016 onwards and exams in 2018 onwards (version 1.1). It covers every point from 4.6.1.1 to 4.6.2.7: calculating rates, the factors that affect rate, collision theory, catalysts, reversible reactions and equilibrium. The topic is assessed on Paper 2 (topics 6 to 10), set at Foundation and Higher Tier. Content the specification marks (HT only) is labelled Higher tier only; that includes all of 4.6.2.4 to 4.6.2.7. Required practical activity 5 belongs to this topic.

For quick recall, use the revision notes. To test yourself, use the practice questions. The course hub is AQA GCSE Chemistry, and the printable checklist lists every specification point. Activation energy and reaction profiles are introduced in the Energy changes study guide.

What this unit covers

Spec point	What you must be able to do	Tier
4.6.1.1 Calculating rates	Calculate mean rate in g/s or cm³/s; draw and interpret graphs; draw tangents	Both
4.6.1.1 Calculating rates	Use moles and mol/s; calculate the gradient of a tangent	Higher tier only
4.6.1.2 Factors	Recall how concentration, pressure, surface area, temperature and catalysts affect rate	Both
Required practical 5	Investigate how concentration affects rate by a gas volume method and a colour or turbidity method	Both
4.6.1.3 Collision theory	Explain rate changes with collision theory; surface area to volume ratio; proportionality	Both
4.6.1.4 Catalysts	Identify catalysts; explain catalytic action using activation energy	Both
4.6.2.1–4.6.2.3	Reversible reactions, energy changes in each direction, equilibrium in a closed system	Both
4.6.2.4–4.6.2.7	Le Chatelier’s Principle; effects of concentration, temperature and pressure	Higher tier only

4.6.1.1 Calculating rates of reactions

You find the rate by measuring the quantity of a reactant used or a product formed over time:

mean rate of reaction = quantity of reactant used ÷ time taken
mean rate of reaction = quantity of product formed ÷ time taken

Quantities are measured as a mass in grams or a volume in cm³, so units are g/s or cm³/s.

Worked example. A reaction makes 60 cm³ of gas in 40 s. Mean rate = 60 ÷ 40 = 1.5 cm³/s.

Worked example. A flask loses 1.20 g as carbon dioxide escapes over 150 s. Mean rate = 1.20 ÷ 150 = 0.008 g/s.

Moles and mol/s (Higher tier only)

Convert mass to moles with moles = mass ÷ Mr (see Quantitative chemistry).

Worked example. 0.88 g of carbon dioxide (Mr = 44) is produced in 40 s.

moles = 0.88 ÷ 44 = 0.020 mol
rate  = 0.020 ÷ 40 = 0.0005 mol/s  (5.0 × 10⁻⁴ mol/s)

Graphs and tangents

Plot quantity of product (or reactant) on the y-axis against time on the x-axis and draw a smooth curve.

• The curve is steepest at the start, when reactant concentration is highest.
• It gets less steep as reactants are used up.
• It becomes flat when the reaction stops.

The slope of a tangent to the curve at a point measures the rate at that moment. A steeper tangent means a faster rate.

Gradient of a tangent (Higher tier only). Draw a straight line that just touches the curve at the chosen time. Pick two points far apart on the tangent and calculate change in y ÷ change in x.

Worked example. A tangent at 25 s passes through (5 s, 12 cm³) and (45 s, 60 cm³). Gradient = (60 − 12) ÷ (45 − 5) = 48 ÷ 40 = 1.2 cm³/s.

4.6.1.2 Factors which affect the rates of chemical reactions

The rate increases when you increase any of:

• the concentration of reactants in solution
• the pressure of reacting gases
• the surface area of solid reactants (smaller pieces or powder)
• the temperature

and when you add a catalyst.

Required practical 5

Investigate how changes in concentration affect the rate by two methods, and develop a hypothesis first (for example “the higher the concentration, the faster the rate”).

Method 1 – volume of gas. React magnesium ribbon with hydrochloric acid of different concentrations. Collect the hydrogen in a gas syringe (or an upturned measuring cylinder in water). Record the volume every 10 s.

Method 2 – colour or turbidity. Mix sodium thiosulfate solution with hydrochloric acid in a flask on a paper cross. The mixture turns cloudy as sulfur forms. Time how long it takes for the cross to disappear.

For the second method, rate is proportional to 1/time.

Worked example. At 0.10 mol/dm³ thiosulfate the cross disappears after 80 s; at 0.20 mol/dm³ after 40 s.

1/t at 0.10 mol/dm³ = 1 ÷ 80 = 0.0125 s⁻¹
1/t at 0.20 mol/dm³ = 1 ÷ 40 = 0.025 s⁻¹

Doubling the concentration doubles 1/t, so the rate is proportional to concentration here.

Control variables: temperature, volume of each solution, the same cross and viewer (method 2), the same length of magnesium (method 1).

4.6.1.3 Collision theory and activation energy

Reactions happen only when particles collide with sufficient energy. The minimum energy needed is the activation energy.

Change	Effect on particles	Effect on rate
Higher concentration (solutions)	More particles in the same volume, so more frequent collisions	Increases
Higher pressure (gases)	Particles closer together, so more frequent collisions	Increases
Larger surface area (solids)	More particles exposed, so more frequent collisions	Increases
Higher temperature	Particles move faster, so more frequent collisions and collisions are more energetic – more have at least the activation energy	Increases

Temperature is the only factor with two effects. Say both.

Proportionality. If doubling the concentration doubles the number of particles in the same volume, the frequency of collisions roughly doubles, so the rate roughly doubles.

Surface area to volume ratio

Cutting a solid into smaller pieces keeps its volume but increases its surface area.

Worked example. A 2 cm cube: surface area = 6 × 2 × 2 = 24 cm², volume = 8 cm³, ratio = 24 : 8 = 3 : 1. Cut into eight 1 cm cubes: surface area = 8 × 6 = 48 cm², volume still 8 cm³, ratio = 6 : 1.

The ratio doubles, so more particles are exposed and collisions with the solid are more frequent.

4.6.1.4 Catalysts

• A catalyst changes the rate of a reaction but is not used up.
• Different reactions need different catalysts.
• Enzymes act as catalysts in biological systems.
• A catalyst provides a different pathway with a lower activation energy, so more collisions have enough energy to react.

On a reaction profile for a catalysed reaction, the reactant and product levels are unchanged. Only the hump is lower.

Identifying a catalyst. A substance is a catalyst if it speeds up the reaction, its mass is the same at the end, and it does not appear in the chemical equation. You do not need to learn the names of catalysts beyond those in the specification.

4.6.2.1 Reversible reactions

In some reactions, the products can react to produce the original reactants. The sign ⇌ shows this:

A + B ⇌ C + D

The direction can be changed by changing the conditions. For example, heating ammonium chloride makes ammonia and hydrogen chloride; on cooling they react to form ammonium chloride again:

ammonium chloride ⇌ ammonia + hydrogen chloride

4.6.2.2 Energy changes and reversible reactions

If a reversible reaction is exothermic in one direction, it is endothermic in the other. The same amount of energy is transferred in each direction.

hydrated copper sulfate (blue) ⇌ anhydrous copper sulfate (white) + water

Heating blue hydrated copper sulfate is endothermic. Adding water to white anhydrous copper sulfate turns it blue again and is exothermic – the mixture warms up.

4.6.2.3 Equilibrium

When a reversible reaction happens in apparatus that prevents the escape of reactants and products (a closed system), equilibrium is reached when the forward and reverse reactions occur at exactly the same rate. The reactions have not stopped; the amounts stay constant because they happen at the same rate.

4.6.2.4–4.6.2.7 Changing conditions at equilibrium (Higher tier only)

The relative amounts of reactants and products at equilibrium depend on the conditions. Le Chatelier’s Principle: if a system is at equilibrium and a condition is changed, the system responds to counteract the change. Questions give you the information you need; you make qualitative predictions.

Concentration (4.6.2.5)

If a concentration changes, the system is no longer at equilibrium, and all concentrations change until equilibrium is reached again.

• Increase the concentration of a reactant → more products form.
• Decrease the concentration of a product → more reactants react.

Temperature (4.6.2.6)

Temperature	Forward reaction endothermic	Forward reaction exothermic
Increased	Relative amount of products increases	Relative amount of products decreases
Decreased	Relative amount of products decreases	Relative amount of products increases

Pressure (4.6.2.7) – gaseous reactions only

• Increased pressure → equilibrium shifts towards the side with fewer molecules.
• Decreased pressure → equilibrium shifts towards the side with more molecules.

Count molecules from the balanced symbol equation.

Worked example

2SO₂(g) + O₂(g) ⇌ 2SO₃(g). The forward reaction is exothermic.

• Pressure increased: left side has 2 + 1 = 3 molecules; right side has 2. The equilibrium shifts right, so the yield of SO₃ increases.
• Temperature increased: the forward reaction is exothermic, so the relative amount of SO₃ decreases.
• More O₂ added: a reactant’s concentration increases, so more SO₃ forms.

If both sides have the same number of gas molecules, for example H₂ + I₂ ⇌ 2HI, changing the pressure has no effect on the position of equilibrium.

Common errors

• Giving a rate without units, or using cm³ when the data are in grams.
• Drawing a tangent that cuts through the curve.
• Saying particles collide “more” – write “more frequent collisions” or “more collisions per second”.
• For temperature, giving only the collision frequency point and missing “more energetic collisions”.
• Saying a catalyst “gives particles more energy”. It lowers the activation energy.
• Saying the reactions stop at equilibrium.
• Using the pressure rule on a reaction with no gases, or counting atoms instead of molecules.

Where to go next

• Recall and self-test: revision notes
• Exam-style practice: practice questions
• Find your weak spots: free 10-minute diagnostics

Official syllabus

AQA GCSE Chemistry (8462) specification, for teaching from September 2016 onwards, exams in 2018 onwards, version 1.1 (4 October 2019), published by AQA – section 4.6 The rate and extent of chemical change.